Kinetics by the Initial Rates Method
Introduction
Rate laws are mathematical expressions that describe the relationship between reactant concentrations and the rates of reaction, taking the form,
Rate = k[reactant 1]m[reactant 2] n[reactant 3]p… [M/s] (1)
The proportionality constant k is the rate constant; any given reaction has a specific value of k for a given set of conditions, such as temperature, pressure, and solvent; the rate constant, in contrast to the rate of reaction, does not depend on the concentration of the reactants. Exponents m, n, p, …, are the reaction orders, and indicate the degree to which the reaction rate depends on the concentration of the associated reactant. “There is no simple correlation between the stoichiometry of the reaction and the rate law”; the reaction orders, and thus the rate constant, must be determined experimentally, as in this experiment between acetone (CH3COCH3), the hydronium ion (H3O+), and iodine (I2), called the iodination of acetone. The reaction in this experiment is shown in Figure 1 and Eq.2.2 Figure 1. Reaction of acetone, iodine, and hydronium ion.
CH3COCH3 (aq) + I2 (aq) CH3COCH2I (aq) + HI (aq) (2)
Thus, the rate law in this experiment is shown in Eq. 3.
Rate = k[CH3COCH3]m[H3O+] n[I2]p (3)
The reaction rate expressions are shown in Eq. 4.
Rate = -Δ[CH3COCH3]/Δt = -Δ[H3O+]/Δt = -Δ[I2]/Δt (4)
Spectrophotometry can be used here to measure the reaction rate expressed as the disappearance of I2, since I2 absorbs in the visible light spectrum; initially violet, the color decreases in intensity as I2 is consumed, and the measured absorbance decreases as more light reaches the spectrophotometer detector. By Beer’s Law, A = εbc, absorbance A is directly proportional to the concentration c of the absorbing substance, I2, (when working with changes in absorbance, molar absorptivity